Changes and Properties of Matter

1. Changes and Properties of Matter

2. Physical Properties of Matter • Physical Changes: Changes that change only the appearance of a substance, not its chemical identity. • Physical Properties: Properties that can be observed through physical change

3. Physical Change Physical Property Melting point (the temperature at which a substance turns from a solid to a liquid) Boiling point (the temperature at which a substance turns from a liquid to a gas) Solubility (the amount of solute that can be dissolved in a solvent) Vapor pressure (the pressure exerted by a vapor at vapor-liquid equilibrium) Malleability (the ability to be hammered or rolled into thin sheets) Ductility (the ability to be stretched into a wire) Density (the mass of a substance per unit volume) Electrical & Heat Conductivity (the ability to pass heat or electricity through a substance • Melting • Boiling • Dissolving • Evaporating • Crushing • Stretching

4. Chemical Properties and Changes • Chemical Changes: Changes that result in changing the chemical composition of a substance. Can be reversed only by another chemical change. • Chemical Properties: Properties that can only be observed through chemical change.

5. Chemical Change Chemical Property Reactivity (the likelihood of one substance to undergo a chemical reaction with another substance) Stability (the likelihood that a substance will no decompose) Heat of Reaction (the energy absorbed or released by a chemical reaction) • Corrosion of metals, flammability • Chemical decomposition e.g. hydrogen peroxide decomposes to form water and oxygen, but water does not decompose spontaneously • Combustion releases heat (exothermic) Rarely do chemical changes absorb heat (endothermic)

6. Indications of a chemical reaction: • Energy absorbed or released (heat and/or light) • Color change • Gas production (bubbling, fizzing, or odor change) • formation of aprecipitate- a solid that separates from solution (won’t dissolve) • Irreversibility- not easily reversed • new substances produced, old substances destroyed.

7. Checking for understanding Explain why are these processes physical changes: • Boiling • Condensation • Shredding paper • Explain why are these processes chemical changes: • Burning paper • Rusting bicycle • Frying eggs

8. Change in Energy

9. Energy Different types of energy exist (heat, potential, kinetic, chemical, nuclear etc.) Heat - the energy transferred between objects that are at different temperatures. Unit of heat is joules J. According to the Law of Conservation of Energy, energy cannot be created or destroyed. It can only by transferred or transformed. The total amount of energy in a system remains constant.

10. Add a potential energy diagram • To draw from scratch • Do more calculations on enthalpy/energies

11. Potential Energy Diagram http://www.kentchemistry.com/links/Kinetics/PEDiagrams.htm

12. Potential Energy Diagram • Potential Energy Diagrams show changes in energy during chemical reactions. • CHEMICAL EQUATIONS A + B ---> C + D ReactantProducts • In chemical reactions, reactants react to form products with an associated release or absorption of energy.

13. Change in Heat Energy • Exothermic Reaction • The change in heat (∆H) is NEGATIVE. • Heat is released into the surroundings • Endothermic Reaction • The change in heat (∆H) is POSITIVE. • Heat is absorbed by the system

14. Potential Energy Diagram Potential Energy in KJ N Energy O O N CO + NO2 C C C O O O REACTANTS O O N O O No unit, fancy of saying “start to finish” PRODUCTS Reaction Coordinate

15. Potential Energy Diagram Heat of Activated Complex Energy CO + NO2 C Heat of Reactants REACTANTS A PRODUCTS Heat of Products D Reaction Coordinate

16. Potential Energy Diagram Activation Energy Amount of energy needed to start the reaction N Energy O O B N CO + NO2 C C C O O O REACTANTS O O N ΔH O O ΔH: ENTHALPY Net energy of the reaction PRODUCTS Reaction Coordinate

17. Enthalpy (H)is used to quantify the heat flow into or out of a system in a process that occurs at constant pressure. DH = H (products) – H (reactants) DH = heat given off or absorbed during a reaction at constant pressure Hproducts< Hreactants Hproducts > Hreactants DH < 0 DH > 0

18. Activation Energy with no catalyst Potential Energy→ Activation Energy with a catalyst Enthalpy of Reaction Reaction Coordinate Potential Energy Diagrams • A catalyst is a substance that changes the speed of reactions by changing the activation energy. • Catalysts speed up chemical reaction by lowering the activation energy • Catalysts are not consumed during the reaction and can be recovered

19. Click Below for the Video Lectures Activation Energy Reaction Coordinate Catalyst

20. Exothermic and Endothermic

21. Exothermic Reaction A + B C + D + ENERGY • Heat is released as a product. • This heat is released into the environment surrounding the reaction, causing the temperature of the surroundings to increase. • The products have less energy than the reactants after release. • Since the products have less energy stored in their chemical bonds, they are more stable than the reactants were. Burning paper is exothermic. The ash formed by the burning is not flammable and will not burn.

22. General Exothermic Reaction A + B C + 70 kJ ΔHA = 60 kJ ΔHB = 40 kJ ΔHc = 30 kJ Energy Released as product Exothermic 100 KJ Law of conservation of energy states that the total energy of an isolated system cannot change.

23. Exothermic Reaction C (g) + O2 (g) CO2 (g) ΔH = -393.5 kJ • The minus sign in front of the enthalpy 393.5 kJ indicates that this reaction is exothermic, that energy was released. • This energy can be placed on the products side, as follows: C (g) + O2 (g)  CO2 (g) + 393.5 kJ

24. Synthesis Decomposition C (g) + O2 (g) CO2 (g) ΔH = -393.5 kJ • If we were to synthesize 2.3 moles of CO2, how many kJ would be released? 1 mol of CO2 2.3 mol CO2 393.5 kJ = - 910 kJ 1 mol CO2 If ΔHsynthesis= - 393.5 kJ/mole, then Δ Hdecomposition= +393.5 kJ/mole

25. Potential Energy Diagram - Exothermic A + B C + D + 40 kJ 150 Activation Energy = 150 – 70 80 KJ Energy ΔH = 70 – 30 - 40 kJ, released A + B 70 REACTANTS C + D 30 PRODUCTS Reaction Coordinate

26. Endothermic Reaction A + B + ENERGYC + D • Heat is absorbed by the reactants. • This heat is absorbed from the environment surrounding the reaction, and the temperature of the surroundings decreases. • The products have more energy than the reactants after absorption. • This energy is stored in the chemical bonds of the products

27. General Endothermic Reaction A + B C + 50 kJ ΔHA = 40 kJ ΔHB = 20 kJ ΔHc = 110 kJ Energy absorbed as product Endothermic 60 KJ Law of conservation of energy states that the total energy of an isolated system cannot change.

28. Potential Energy Diagram - Endothermic A + B C + D + 40 kJ 150 Energy C + D 70 PRODUCTS ΔH = 70 – 30 + 40 kJ, Absorbed Activation Energy = 150 – 30 80 KJ A + B 30 REACTANTS Reaction Coordinate

29. Click Below for the Video Lectures Exothermic and Endothermic

30. Calorimetry

31. Calorimetry q = mCΔT Joules (J) Grams (g)J/goCoC units m is the mass of water in the calorimeter cup that absorbs heat from the change or releases heat to the change. q is the quantity of heat that is absorbed or released by a physical or chemical change C is the specific heat of water, the rate at which water gains or loses heat if energy is absorbed or removed from it. ∆T is the Temperature change of the water in the calorimeter cup as a result of the physical or chemical change. What Each Variable Means

32. CSpecific Heat The amount of heat required to raise the temperature of one gram of substance by one degree Celsius. cal/g°C J/g°C (or J/gK) water 1.00 4.18 aluminum0.22 0.90 copper 0.093 0.39 silver 0.057 0.24 gold 0.031 0.13

33. Use this equation on the back of periodic table Specific Heat of water. Use the one for correct unit! They are NOT equations!!!

34. Circle the numbers, underline what you are looking for. • Make a list of number you circled using variables. • Write down the formula • Derive the formula to isolate the variable you are looking for. • Plug in the numbers • How many joules are absorbed by 100.0 grams of water if the temperature is increased from 35.0oC to 50.0oC? q = mcΔT m = 100.0 g 35.0oC 50.0oC ? 4.18 j/goC T i= q = (100.0 g)(4.18 j/goC)(50.0oC- 35.0oC) T f= q = 6270 J q= C=

35. Circle the numbers, underline what you are looking for. • Make a list of number you circled using variables. • Write down the formula • Derive the formula to isolate the variable you are looking for. • Plug in the numbers • 300. J of energy is absorbed by a 50. g sample of water in a calorimeter. How much will the temperature change by? q = mcΔT q = 300 J 50g ? 4.18 j/goC ΔT = q / mc m = ΔT = ΔT = 300 J/ (50g)(4.18 j/goC) C= ΔT = = 1.4354066 oC ΔT = = 1.4oC

36. Checking for understanding • How many joules are required to raise the temperature of 100. grams of water from 20.0oC to 30.0oC ? • How many grams of water can be heated from 20.0oC to 75.0oC using 2500. J?

37. Calorimetry • We use an insulated device called a calorimeter to measure this heattransfer. • A typical device is a “coffeecup calorimeter.”

38. Calorimetry To measure ΔH for a reaction • dissolve the reacting chemicals in known volumes of water • measure the initial temperatures of the solutions • mix the solutions • measure the final temperature of the mixed solution

39. Calorimetry • The heat generated by the reactants is absorbed by the water. • We know the mass of the water, mwater. • We know the change in temperature, ∆Twater. • We also know that water has a specific heat of Cwater = 4.18 J/°C-g. • We can calculate the heat of reaction by: • qsys = ∆H = −qsurr = -mwater × Cwater × ∆Twater

40. When 25.0 mL of water containing 0.025 mol of HCl at 25.0°C is added to 25.0 mL of water containing 0.025 mol of NaOH at 25.0°C in a coffee cup calorimeter, a reaction occurs. Calculate ∆H (in kJ) during this reaction if the highest temperature observed is 32.0°C. Assume the densities of the solutions are 1.00 g/mL. Knowns: Vfinal = VHCl + VNaOH = (25.0 + 25.0) mL = 50.0 mL Dwater = 1.00 g/mL ∆Twater = Tfinal − Tinitial = 32.0°C − 25.0°C = +7.0°C Cwater = 4.18 J/°C-g Calculation: mwater = 50.0 g ∆H = −m × C × ∆T = −(50.0 g)(4.18 J/°C-g)(7.0°C) = −1463 J = −1.5×103 J

41. Phase Change Heating and Cooling Curve

42. Condense Freeze Evaporate Melt Sublime Gas Liquid Solid Deposit

43. Phase Change • During a phase change, there is a change in heat energy but no change in temperature. At the melting point, heat energy is being used to break down the crystalline lattice. At the boiling point, heat energy is being used to convert the liquid to a gas. During a phase change, heat is being used to change phase, not increase the kinetic energy of the particles. It is for this reason that Temperature remains constant in a phase change while heat energy changes.

44. LIQUID TO GAS GAS EVAPORATING BP D E Temperature SOLID to LIQUID LIQUID B MELTING MP C SOLID A Energy

45. Phase Change • During phase change, temperature stays constant but energy increases.

46. Phase Change • Energy needed to melt is called Latent Heat of Fusion ( ΔHfus) • Energy need to vaporize is called Latent Heat of Vaporization (ΔHvaporization)

49. Triple point – the point on a phase diagram at which the three states of matter: gas, liquid, and solid coexist Critical point – the point on a phase diagram at which the substance is indistinguishable between liquid and gaseous states

50. q = mCΔT q = mCΔT q = mCΔT